Direction of Reactions

Effects of a Side Chain at the 3 Position on the Direction of Reaction Pathway.. 134... 

By the mid-1930s, catalytic technology entered into petroleum refining. To a greater extent than thermal cracking, catalysis permitted the close control of the rate and direction of reaction. It minimized the formation of unwanted side reactions, such as carbon formation, and overall improved the yield and quality of fuel output. 

Strategy First calculate the partial pressures of N204 and N02, using the ideal gas law as applied to mixtures P, = tiiRT/V. Then calculate Q. Finally, compare Q and K to predict the direction of reaction. 

We have seen that the value of an equilibrium constant tells us whether we can expect a high or low concentration of product at equilibrium. The constant also allows us to predict the spontaneous direction of reaction in a reaction mixture of any composition. In the following three sections, we see how to express the equilibrium constant in terms of molar concentrations of gases as well as partial pressures and how to predict the equilibrium composition of a reaction mixture, given the value of the equilibrium constant for the reaction. Such information is critical to the success of many industrial processes and is fundamental to the discussion of acids and bases in the following chapters. 

The reactants appear on the left and the products appear on the right. The arrow joining them indicates the direction of reaction. 

The sodium of the organosodium compound is less effective as a Lewis acid than is the magnesium of the Grignard reagent, which may account for some of the differences. With sodium as the metal, the direction of reaction seems to be decided by steric effects, the bulkier reagents avoiding the secondary position. 

The fruitful relationship between experiment and theory has pushed carbene chemistry further toward the direction of reaction control that is, regio- and stereoselectivity in intra- and intermolecular addition and insertion reactions. The interplay between experiment and modem spectroscopy has led to the characterization of many carbenes that are crucial to both an understanding and further development of this held. 

In each case, we are considering only the direction of reaction indicated. The reverse reaction may well be of a different order for example, the decomposition of NO, is second-order.)... 

If we wished to be wholly consistent, we could write both reactions as reduction processes. Reversing the direction of reaction (2) means that we need to change the sign of its contribution toward the overall value of AGr, so... 

Dashed line is the minimum energy path and the arrows indicate the direction of reaction. 

The use of PheDH mutants in the opposite direction of reaction to accomplish the quantitative removal of the L-enantiomer of a range of substituted phenylalanine derivatives is discussed in more detail in Section 5.03.3.4. 

Here kx is the temperature dependent rate constant for the forward direction of reaction step 1, which is assumed to follow an Arrhenius expression with activation energy of Ej of Figure 4.33, is the pressure of the reactant A2, 0t is the fraction... 

As explained in Section 6.3.11, the inner potential difference—A( )—seems to encompass all the sources of potential differences across an electrified interface—Ax and A jf—and therefore it can be considered as a total (or absolute ) potential across the electrode/electrolyte interface. However, is the inner potential apractical potential First, the inner potential cannot be experimentally measured (Section 6.3.11). Second, its zero point or reference state is an electron at rest at infinite separation from all charges (Sections 6.3.6 and 6.3.8), a reference state impossible to reach experimentally. Third, it involves the electrostatic potential within the interior of the phase relative to the uncharged infinity, but it does not include any term describing the interactions of the electron when it is inside the conducting electrode. Thus, going back to the question posed before, the inner potential can be considered as a kind of absolute potential, but it is not useful in practical experiments. Separation of its components, A% and A f, helped in understanding the nature of the potential drop across the metal/solution interface, but it failed when we tried to measure it and use it to predict, for example, the direction of reactions. Does this mean then that the electrochemist is defeated and unable to obtain absolute potentials of electrodes ... 

Equation 9.72 introduces a great deal of nomenclature at once. Chemical species are indexed by k, with K being the total number of species (later, when we generalize the kinetics to multiple phases, the variable Kg is used for the number of gas-phase species) reactions are indexed by the variable i, with / being the total number of reactions in the mechanism the name of species k is represented by X v ki is the stoichiometric coefficient of species k in the forward direction of reaction i is the stoichiometric coefficient of species k in the reverse direction of reaction i. 

A and B is converted to the internal (vibrational) energy of the association product. By analogy, the highly excited species formed is denoted C. Because it is highly unstable, C may unimolecularly dissociate by the forward direction of reaction 9.101, with rate constant kd, or if C collides with another species M, it could be stabilized via the reverse of reaction 9.100, forming the product C. 

Within each pool, the direction of reaction is determined by simple mass-action considerations. The direction of overall conversion (glycolysis or gluconeogenesis) depends on the rates at which the connecting reactions supply and remove materials from the pool. 

The reaction quotient Qc is useful because it lets us predict the direction of reaction by comparing the values of Qc and Kc. If Qc is less than Kc, movement toward equilibrium increases Qc by converting reactants to products (that is, net reaction proceeds from left to right). If Qc is greater than Kc, movement toward equilibrium decreases Qc by converting products to reactants (that is, net reaction proceeds from right to left). If Qc equals Kc, the reaction mixture is already at equilibrium, and no net reaction occurs. 

The value of the equilibrium constant for a reaction makes it possible to judge the extent of reaction, predict the direction of reaction, and calculate equilibrium concentrations (or partial pressures) from initial concentrations (or partial pressures). The farther the reaction proceeds toward completion, the larger the value of Kc. The direction of a reaction not at equilibrium depends on the relative values of Kc and the reaction quotient Qc, which is defined in the same way as Kc except that the concentrations in the equilibrium constant expression are not necessarily equilibrium concentrations. If Qc Kcr net reaction goes from left to right to attain equilibrium if Qc > Kc/ net reaction goes from right to left if Qc = Kc/ the system is at equilibrium. 

For each of the following equilibria, use Le Chatelier s principle to predict the direction of reaction when the volume is increased. 

To predict the direction of reaction, use the balanced equation to identify the proton donors (acids) and proton acceptors (bases), and then use Table 15.1 to identify the stronger acid and the stronger base. When equal concentrations of reactants and products are present, proton transfer always occurs from the stronger acid to the stronger base. 

---